Brønsted-Lowry and Lewis definitions, pH calculations, buffers, and titration curves - step-by-step worked examples for H2 Chemistry.
Q: What does H2 Chemistry Notes: Topic 4 - Theories of Acids and Bases cover? A: Compare Arrhenius, Brønsted-Lowry, and Lewis definitions while mastering pH calculations, buffer design, and titration curves for Core Idea 2 (Theories of Acids and Bases).
Acid-base theory appears across Papers 1-4, from conceptual MCQs to practical titrations. This note organises the definitions, equilibrium expressions, and exam-standard calculations required in 2026. Continue exploring linked drills at https://eclatinstitute.sg/blog/h2-chemistry-notes.
Status: SEAB's current H2 Chemistry (9476) syllabus PDF is labelled for 2026, and the current Chemistry Data Booklet is labelled 8873/9476/9813 for use from 2026 in non-practical papers. Core Idea 2 Topic 4 scope is assessed across Papers 1-3.
The core idea is simple: Acid-base questions ask what is transferred: H+, OH-, or an electron pair.
Use it as a working check: Choose the model first, then calculate. Bronsted-Lowry is usually about proton transfer; Lewis is usually about electron-pair donation and acceptance.
Then go one layer deeper: Example: ammonia is a Bronsted base when it accepts H+ to form ammonium, and a Lewis base when its lone pair bonds to a metal ion.
Quick revision box
What this topic tests: Acid-base models, pH/buffer calculations, and titration curve interpretation.
Top mistakes to avoid: Using wrong acid-base model; buffer equation misuse; poor link between curve features and chemistry.
20-minute sprint plan: 5 min Arrhenius/Brønsted/Lewis contrast; 10 min pH-buffer-titration drill; 5 min curve annotation.
Route map: choose the acid-base tool first
If the question asks about...
Start with...
Then connect to...
Trap to avoid
Definition or role of a species
The named acid-base model
Proton transfer for Brønsted-Lowry, electron-pair donation or acceptance for Lewis
Check this topic from memory
Attempt the matching topic bank before reopening the notes. Use each missed idea to decide what to review next.
Do not mix models in the same sentence unless the question asks for comparison.
Strong acid or strong base pH
Complete dissociation
Concentration of H+ or OH-, then pH or pOH
Do not use weak-acid equilibrium tables for a fully dissociated species.
Weak acid or weak base pH
Equilibrium setup
Initial concentration, change, equilibrium amount, then the acid or base constant
Do not assume the approximation without checking whether it is reasonable.
Buffer after acid or base is added
Mole adjustment first
Remaining conjugate pair ratio, then Henderson-Hasselbalch
Do not put initial concentrations straight into the buffer equation after reaction.
Titration-curve feature
Stage of the titration
Initial pH, buffer region, half-neutralisation, equivalence, or excess reagent
Do not call every equivalence point pH 7.
Indicator choice
Equivalence-region pH
Whether the indicator transition range overlaps the sharp pH jump
Do not choose an indicator from memory if the paper supplies a range table.
Use this route map before calculating. Acid-base questions often lose marks because the calculation is correct for the wrong stage or the explanation uses the wrong model.
State the theory explicitly in answers. For example, “According to Brønsted-Lowry theory, ammonia acts as a base because it accepts a proton to form NHX4X+.”
Lewis donor-acceptor checkpoint
For Lewis acid-base questions, mark the electron-pair donor and acceptor before writing the bond formed. This prevents a common error: naming the proton-transfer role when the question is really testing lone-pair donation to an electron-deficient centre.
Example
Lewis base
Lewis acid
Bond-forming idea
Common trap
NHX3+HX+NHX4X+
NHX3, using the nitrogen lone pair.
HX+, accepting the electron pair.
The donated lone pair forms the N−H bond.
Saying HX+ donates a proton, then forgetting the electron-pair view.
Worked check: in BFX3, boron has only six electrons around it. When NHX3 approaches, the nitrogen lone pair is donated to boron, so NHX3 is the Lewis base and BFX3 is the Lewis acid.
Misconception check: a Lewis acid does not need to contain hydrogen. It only needs to accept an electron pair.
Definitions above follow the IUPAC Gold Book. The SEAB Chemistry Data Booklet provides Kw (at 298K) and key constants; Ka/Kb values are provided in the question when needed.
2 Quantifying Acid Strength
2.1 Ka and pKa
For a weak acid HA, the acid dissociation constant is
Ka=[HA][HX+][AX−],pKa=−log10Ka.
Smaller pKa values indicate stronger acids (greater dissociation). Always include state symbols when writing equilibrium equations, for example:
HA(aq)+HX2O(l)HX3OX+(aq)+AX−(aq).
2.2 Polyprotic Acids
Each deprotonation has its own Ka. When calculating pH, consider whether steps are independent. Typically, the first dissociation dominates, and subsequent Ka values are much smaller.
3 pH Calculations
3.1 Strong Acids and Bases
Assume complete dissociation:
Solute
Reaction
pH/pOH shortcut
HCl
HClHX++ClX−
pH=−log10[HX+]
NaOH
NaOHNaX++OHX−
pH+pOH=14 assumes Kw=1.0×10−14 at 298K from the SEAB Chemistry Data Booklet; if the paper supplies a different temperature or Kw, adjust accordingly.
3.2 Weak Acids/Bases
Use equilibrium tables. Example for CH3COOH:
Species
Initial
Change
Equilibrium
CHX3COOH
C
−x
C−x
HX+
0
+x
x
CHX3COOX−
0
Then Ka=C−xx2. If x is much smaller than C, you can approximate C−x≈C to simplify the algebra.
Weak-acid approximation checkpoint
Before using C−x≈C, check that the assumption is reasonable. If the calculated x is not small compared with C, keep C−x and solve the quadratic.
Situation
First setup
Approximation check
Common trap
Ka is much smaller than C
Ka=C−xx2
Try x≈KaC, then check Cx≤0.05
Treating the approximation as automatic for every weak acid.
Cx is more than about 5
Keep C−x.
Solve x2+Kax−KaC=0
The weak acid is very dilute
Start from the full expression.
Use the equation supplied by the question or solve more carefully.
Ignoring that dilution can make the shortcut less reliable.
Worked check: for 0.100molL−1 ethanoic acid with Ka=1.8×10−5, x≈(1.8×10−5)(0.100)=1.34×10−3. Then Cx=0.0134=1.34, so the approximation is acceptable.
Misconception to avoid: weak acid does not mean x is always negligible. The small-x step is a checked shortcut, not a definition.
After solving for x, compute pH: pH=−log10x.
Ethanoic acid: weak-acid reference for Ka, ICE tables, and buffer calculations.
The −COOH group is the acidic site, while the stabilised conjugate base CHX3COOX− supports the weak-acid equilibrium model used in Paper 2 calculations.
Use the Ka or Kb values given in the question rather than memorised numbers (they are not listed in the SEAB Chemistry Data Booklet).
Calculation order checkpoint
For acid-base calculations, decide the stage before choosing the equation. A good default sequence is:
Identify the species present before any reaction.
Use stoichiometry if strong acid or strong base has been added.
Decide what species remain after reaction.
Choose the equilibrium or buffer expression only after the remaining species are known.
Convert the final [HX+], [OHX−], or conjugate-pair ratio into pH.
Stage in the question
First move
Equation or idea
Common trap
Fresh strong acid or base
Convert concentration directly to [HX+] or [OHX−].
Complete dissociation, then pH or pOH.
Setting up an equilibrium table for a species that fully dissociates.
Fresh weak acid or base
Build an initial-change-equilibrium table.
Ka or Kb.
Assuming the small-x approximation without checking whether it is reasonable.
Buffer before addition
Confirm both conjugate species are present.
Henderson-Hasselbalch.
Using the equation when only one member of the conjugate pair is present.
Buffer after strong acid or base is added
Adjust moles by neutralisation first.
Stoichiometry first, then Henderson-Hasselbalch.
Substituting the original concentrations directly into the buffer equation.
Half-neutralisation point
Show that weak acid and conjugate base amounts are equal.
pH=pKa for a weak acid titrated by strong base.
Quoting the shortcut without proving the stage.
Equivalence point
Identify the salt or excess reagent left in solution.
Hydrolysis or excess strong reagent, depending on the titration.
Assuming every equivalence point has pH=7.
Misconception check: Henderson-Hasselbalch is not the first step after strong acid or strong base is added to a buffer. The neutralisation mole table comes first; the buffer equation comes only after the new conjugate-pair amounts are known.
4 Buffers and Henderson-Hasselbalch
Buffers resist pH change when small amounts of acid or base are added. For a weak acid buffer:
pH=pKa+log10([HA][AX−])
Use the Henderson-Hasselbalch equation only after confirming both species are present in appreciable amounts. When strong acid/base is added, adjust moles first, then recompute concentrations before applying the equation.
Design principle: Choose pKa close to the desired buffer pH (within ±1) for maximum capacity.
For basic buffers, the same logic applies to the NHX4X+/NHX3 pair: NHX3 uses its nitrogen lone pair to accept HX+, reducing sudden pH swings.
Ammonia: Brønsted-Lowry base and common weak-base buffer component.
5 Titration Curves
Recognise four canonical curves: strong acid vs strong base, strong acid vs weak base, weak acid vs strong base, weak acid vs weak base. Key checkpoints:
Initial pH: Based on identity of the analyte.
Buffering region: Exists when a weak acid/base titrated with strong counterpart.
Half-neutralisation point:pH=pKa (or pOH=pKb).
Equivalence point: Determine species present; e.g. weak acid + strong base yields basic salt solution.
Titration-curve stage checkpoint
Before reading a titration curve, name the stage by volume added. The same curve can test different chemistry at different points.
Stage on the curve
Dominant species to name
What controls pH
Common trap
Before any titrant is added
Original acid or base
Initial strong-species concentration or weak-species equilibrium
Starting with the equivalence-point species.
Before equivalence for weak acid plus strong base
Weak acid and conjugate base mixture
Buffer ratio after neutralisation moles are adjusted
Using initial concentrations instead of post-reaction moles.
Half-neutralisation
Equal weak acid and conjugate base amounts
pH=pKa for weak acid systems
Quoting the shortcut when the amounts are not equal.
Equivalence
Salt solution or neutral strong-acid/strong-base mixture
Hydrolysis for weak acid/strong base or weak base/strong acid; neutralisation for strong/strong
Assuming every equivalence point is pH 7.
After equivalence
Excess strong acid or strong base
Excess reagent moles over total volume
Continuing to use the buffer equation after one conjugate species is exhausted.
Worked check: in ethanoic acid titrated with sodium hydroxide, the half-neutralisation point is controlled by equal CHX3COOH and CHX3COOX− amounts, while the equivalence point is controlled by hydrolysis of CHX3COOX−, so the pH is above 7.
Misconception check: the steep jump is not the same as equivalence. Equivalence is the stoichiometric point; the steep jump is the curve region around it.
State pH values qualitatively if data not supplied (e.g. “equivalence point pH>7 due to hydrolysis of the conjugate base”).
6 Worked Example
Question:25.0mL of 0.100molL−1 ethanoic acid is titrated with 0.100mol⋅L−1 sodium hydroxide. Ka(CHX3COOH)=1.8×10−5. Calculate the pH after 12.5mL of NaOH has been added.
Solution:
Initial moles: n(CHX3COOH)=CV=(0.100molL−1)(25.0mL)=0.00250mol. Added base moles: n(NaOH)=(0.100molL−1)(12.5mL)=0.00125mol.
Statement: pH=4.74 (rounded to two decimal places).
At half-neutralisation, [AX−]=[HA], so pH=pKa - a result worth memorising.
7 Practical Considerations
Indicators: Choose based on equivalence point pH. For weak acid-strong base titrations, phenolphthalein matches the basic equivalence region.
Use indicator ranges provided in the paper; if a table is supplied, cite the range directly instead of relying on memory.
Paper 4 tasks: When planning buffer preparation, describe stepwise calculations (moles before diluting, then the dilution formula C1V1=C2V2), and specify precise volumetric equipment (pipette, burette).
8 Common Pitfalls
Applying Henderson-Hasselbalch without confirming both conjugate pairs remain.
Forgetting to include the auto-ionisation of water for highly dilute strong acids or bases (e.g. when [HX+]<1×10X−6mol⋅L-1).
Mixing units (molarity vs mole).
Assuming equivalence point pH=7 for all titrations.
9 Drill Suggestions
Rank HClO, HClOX2, HClOX3, and HClOX4 by acid strength; justify using inductive effect and oxidation state.
Design a buffer at pH=9.2 using the NHX4X+/NHX3
Sketch titration curves for (a) HNOX3 vs KOH and (b) CHX3COOH
Common exam mistakes
Assuming equivalence point pH is always 7: This is only true for strong acid-strong base titrations. A weak acid-strong base equivalence point gives pH > 7 due to conjugate base hydrolysis; weak base-strong acid gives pH < 7.
Applying Henderson-Hasselbalch before confirming buffer conditions: The equation is only valid when both the weak acid and its conjugate base are present in appreciable amounts. Using it at or near the equivalence point gives a meaningless result.
Forgetting to adjust moles before using Henderson-Hasselbalch: When strong acid or base is added to a buffer, convert moles and recompute the ratio of conjugate pair first; skipping this step produces a wrong pH.
Mixing up conjugate acid-base pairs: The conjugate base of a weak acid is the species formed after proton donation; identifying the wrong species loses marks in Brønsted-Lowry definition questions.
Not specifying the acid-base model in answers: Writing "acts as a base" without stating "according to Brønsted-Lowry theory" is incomplete; the question usually specifies which model to use.
Ignoring water autoionisation at very low concentrations: For highly dilute strong acid solutions where [HX+]<1⋅10−6mol⋅L−1, water's contribution to [HX+] becomes significant and cannot be ignored.
Choosing an indicator with the wrong pH range: The indicator's transition range must overlap with the equivalence point pH. Methyl orange (range ~3-4.5) is unsuitable for weak acid-strong base titrations.
Frequently asked questions
What is the difference between the Brønsted-Lowry and Lewis definitions of a base? A Brønsted-Lowry base is a proton acceptor, while a Lewis base is an electron-pair donor. All Brønsted-Lowry bases are Lewis bases (lone pair donated to accept H⁺), but Lewis bases also include species that donate electron pairs to non-proton acceptors such as metal ions or BFX3. The Lewis definition is broader and covers complex formation and coordination chemistry.
How do I know which acid-base definition to use in a question? Apply the definition that the question specifies or the one that fits the chemistry described. If water is the solvent and proton transfer is occurring, Brønsted-Lowry is appropriate. If the question involves electron pair donation to a metal ion or a Lewis acid like BFX3, use the Lewis definition.
What does pH = pKa at the half-neutralisation point mean in practice? At half-neutralisation, exactly half the weak acid has been converted to its conjugate base, so [HA]=[AX−] and the log term in Henderson-Hasselbalch equals zero. This is a useful shortcut for reading pKa directly from a titration curve at the midpoint volume.
Why does the pH change slowly in the buffering region of a titration curve? In the buffering region, both the weak acid and its conjugate base are present in significant amounts. Adding small amounts of strong base converts some HA to AX− but the ratio changes only slightly, so pH changes very little - this is the definition of buffer action.
Struggling with Theories of Acids and Bases? Our H2 Chemistry tuition programme covers this topic with structured practice, Paper 4 practical drills, and worked exam solutions.
